The
International School of Panama
IB
Chemistry HL a
To
understand how the periodic table was created and how it can be used allows you
to predict:
“the
properties of elements even if you never saw them”.
Before
the development of the modern atomic theory, many elements were discovered. To
be able to understand and predict their properties it was necessary to organize
them.
In
1869, Dimitri Mendeleev published his periodic table.
In
his table the elements were arranged in order of increasing
atomic mass.
The properties of the elements arranged in this way repeated periodically in vertical columns.
Mendeleev
published another version of the periodic table in 1872 in which he left
blank spaces for elements that were not known yet. He predicted their
existence and the properties they should have.
These
elements were later discovered and the properties he predicted were very
accurate for the time.
His
work “can be thought of as similar to putting together a large puzzle.”
(Heath Chemistry)
The
modern periodic table is similar to Mendeleev’s periodic table, but with a
difference proposed by Henry Moseley to solve some discrepancies between
some elements (Ar and K , I and Te). If they are put in order of
increasing atomic mass their properties do not match with those of the elements
in the same column. In
the modern periodic table the elements are in order of increasing atomic
number instead of mass.
PERIODIC LAW: “The properties of the elements repeat periodically when they are arranged in increasing order by their atomic numbers.”
Metals
are at the left (red)
Non-metals
at the right, except for H (blue)
Metalloids
are
between metals and non-metals and they have properties of both.
It
is divided into horizontal rows called periods and vertical
columns called groups.
Groups
are also called families
Why?
Because
the elements in a group share properties like the members of a family.
Group
2
(alkaline earth metals) : 2+
Group
13
(boron group): 3+
The
non metals in:
Group
16:
2-
Group
15:
3-
Group
14: 4-
Now
that you have learned the electron configuration of elements, what relationship
do you see between the configuration, the position in the table and the charge
the elements can get
?
The
elements in the same group have the same outer configuration.
They
tend to acquire the same charge
Because
they all tend to become isoelectronic with the noble gases to complete
their outermost s and p orbitals.
They
have their outer level complete. They don’t “want” more electrons and
they don’t “need” to give away electrons.
…
all
the elements “want” to be like the noble gases so they tend to become
isoelectronic with them
Metallic
properties:
•1)
Increase from top to bottom and from right to left.
Most
active metals: Cs & Fr
•So
the activity of the alkali metals increases from top to bottom. However the
activity of the halogens (non metals of group 17) increases from bottom to top.
•The
most active non metal is F
The
atomic
radius gives
a measure of the size of the atom
http://ull.chemistry.uakron.edu/genobc/Chapter_03/
The
atomic radii decrease from left to right
and from bottom to top in the PT.
Because,
as we move from left to right through a period of the PT, the nuclear
charge increases attracting more the electrons that are added to the same
energy level.
2.
The atomic radii increase from top to bottom. Why?
Because
as we move down a group the electrons are added to higher energy
levels (electrons have the same outermost but with higher principal quantum
number), so they are less attracted by the nucleus. Besides that,
as there are more inner electrons, there is more shielding effect,
which also reduces the attraction of the outer electrons by the nucleus.
Is the energy necessary to remove an electron from a neutral gaseous atom.
Can
be represented by the equation:
X(g)+ IE -à X+ + e-
The
smaller the IE, the greater the tendency of the element to form
a positive ion.
http://ull.chemistry.uakron.edu/genobc/Chapter_03/
http://wulff.mit.edu/pt/pert9.html
In general, the IE increases from left to right in the PT.
However some exceptions are observed in each period.
For
example in period 2, when going from Be to B and from N to O, the
IE decreases. Why?
In
the case of Be to B the decrease is due to the filled 2s orbital
which provides, in the case of B, some shielding to the electron in p.
In
the case of N to O the decrease is due to the extra repulsion in
the doubly occupied p orbital.
There
are various factors that affect the size of an ion:
.The nuclear charge
.The
repulsion of electrons
.The
level of energy of the outer electrons
Comparing
the
ion with the parent atom:
•
Positive ions are smaller than the neutral
atom. Why?
Because
they have the same nuclear charge attracting less electrons,
therefore the attraction is stronger.
For
example K+
< K
•
Negative ions are larger than the neutral
atom. Why?
Because
they have the same nuclear charge attracting more electrons,
therefore the attraction is weaker.
For
example F-
>
F
·
In
a group, the ionic size increases from top to bottom. Why?
The
reason is the same as for the atomic radius:
e- added
to higher energy levels and there is more shielding effect.
In a period, it depends on the type of ions,
but
negative ions decrease from left to right and positive ions
also decrease from left to right. Why?
Because
e-
are added to the same level and
Z (nuclear charge) increases
Isoelectronic
ions
are ions with the same number of electrons
The size of isoelectronic
ions decreases as the nuclear charge (Z) increases. Why?
Because there are more
protons to attract the same amount of electrons distributed in the same
levels and with the same shielding effect.
Example:
O2-
> F-
> Na+
> Mg2+
Electron
Affinity
Is
the energy change associated with the addition of an electron to a gaseous atom.
DE
for : X(g) + e -à
X-
(g)
.
It
can be energy released or absorbed. When it is negative it’s
energy released. The greater the tendency of the element to form negative
ions the more negative the electron affinity.
.
Electron affinities usually become more negative from left to
right in a period of the PT.
.
In
going down a group the electron affinities become more positive
because the electrons added are farther from the nucleus. The change is small
though and there are several exceptions.
Electronegativity
is
the ability of an atom in a molecule to attract shared
electrons to itself.
.
The
electronegativity increases from bottom to top and from
left to right.
http://ull.chemistry.uakron.edu/genobc/Chapter_03/
.
Fluorine is the most electronegative element.
.
Cesium and Francium are the least electronegative elements.
Other
properties like melting and boiling points and densities, depend on the
attraction between particles
However we can observe some patterns in the melting points when moving through a period of the PT, the melting point increases until group 14 which includes solids such as C and Si which form giant molecules with covalent bonds between the atoms. Then the melting points decrease drastically as we get to groups 15, 16, 17 and especially the noble gases in group 18. (We will talk more about this when we study interparticle forces)
